1. Field of the Invention
The present invention relates to a process for capturing and sequestering carbon dioxide and sulfur dioxide using a silicate that is treated with an acid and optionally, a base. Specifically, the present invention relates to serpentine, olivine, or calcium silicate materials, including industrial by-products, like waste cement or calcium-rich fly ash.
2. Description of Related Art
Fossil fuels have been the predominant source of energy in the United States for over a century. However, environmental standards regarding emissions of pollutants into the air are becoming more stringent and as a result the existing state of affairs with respect to fossil fuels is undergoing significant transformation in order to accommodate these stricter standards. One of the greatest obstacles to achieving a minimal environmental impact or “zero emissions” is greenhouse gases, especially carbon dioxide emissions from the burning of fossil fuels.
Anthropogenic emissions have increased carbon dioxide concentrations in the atmosphere to at least 30% more than Pre-Industrial Era levels. Furthermore, it is estimated that future global carbon dioxide emissions will increase from about 7.4 GtC (billions of tons of atmospheric carbon) per year, which was the level in 1997, to 26 GtC per year by 2100. Issues related to the anthropogenic carbon dioxide emissions are closely scrutinized by the public, and the demand for increased carbon reductions, in connection with the large scale utilization of fossil fuels, continues to grow.
The management of carbon dioxide emissions can be achieved by (1) increasing the efficiency of energy conversion; (2) using low-carbon or carbon-free energy sources; and (3) capturing and sequestering carbon dioxide emissions. It is generally accepted that the first two alternatives provide only incremental improvements and, therefore, carbon sequestration technologies are needed to approach zero emissions. Accordingly, there is a clear need for technically sound, cost-effective sequestration processes for carbon dioxide emissions.
Mineral carbonation is a process in which magnesium-rich minerals, such as olivine (Mg2SiO4) and serpentine (Mg3Si2O5(OH)4), react with carbon dioxide to form geologically stable mineral carbonates:(Mg, Ca)xSiyOx-2y+xCO2→x(Mg, Ca)CO3+ySiO2 
These processes mimic the naturally occurring weathering of rocks to form stable magnesium and calcium carbonates, and therefore are based on exothermic reactions. For serpentine minerals, the above reaction can be written as follows:CO2+1/3Mg3Si2O5(OH)4→MgCO3+2/3SiO2+2/3H2O+64 kJ/mole
The mineral carbonation approach has several inherent advantages, including a vast natural abundance of raw materials and the permanent and benign storage of CO2 in solid form. In addition, the overall reaction is exothermic and therefore potentially economically viable. The primary drawback to mineral carbonation is the reaction kinetics, because these naturally occurring processes happen very slowly over geological time instead of industrial time scales. Accordingly, direct carbonation, where the minerals react directly with the carbon dioxide, has been abandoned in favor of processes which the minerals are first suspended in an aqueous carrier, because such systems have now been shown to have much faster kinetics than direct carbonation processes. It is theorized, without wishing to be bound by the theory, that the following set of reactions may occur during the reaction of serpentine with carbon dioxide in aqueous conditions:CO2+H2O→H2CO3→H++HCO3−  (Reaction 1)Mg3Si2O5(OH)4+6H+→Mg2++2SiO2+5H2O  (Reaction 2)Mg2++HCO3−→MgCO3+H+  (Reaction 3)
The initial dissolution of carbon dioxide in water forms carbonic acid, which then dissociates to form bicarbonate ions and hydrogen (Reaction 1). The hydrogen ions then hydrolyze the minerals, liberating Mg2+, and form free silica and water (Reaction 2), or alternatively silicilic acid H4SiO4 (not shown in reaction 2). The dissolved ionic species then forms magnesite, MgCO3, as described in Reaction 3.
Previous studies have focused on crushing magnesium-rich minerals, such as serpentine and olivine, to a fine particle size (<37 μm) to promote surface reactions that are known to control most mineral dissolution reactions. These mineral carbonation studies require extensive communition of the raw materials (<37 μm), high partial pressures (>1950 psig), and long reaction times (>6 hours). However, all of these operations are very energy intensive. For example, pilot scale communition tests to grind the serpentine minerals to minus 200 mesh (75 μm) have reported an energy penalty of 11.5 kWh/ton of mineral processed. Previous studies have also reported that preheating the serpentine minerals to about 650° C. significantly increases their carbonation reactivity, probably due to dehydroxylation, increase in the surface area and destabilization of the crystal structure. However, this pretreatment is very energy intensive, requiring around 200 kWh/ton of serpentine, or in other words, a 20% energy penalty for a coal-fired plant.
Alternatively, magnesium can be first extracted in aqueous solution, however, any acid used for this process must be inexpensive and easily recoverable. Hydrochloric acid and acetic acid have been used to dissolve the magnesium and calcium from serpentine and wollastonite (CaSiO3) minerals, respectively. However, both proposed processes present various technical challenges. For example, in the case of the hydrochloric acid extraction, some of the steps involved require energy input. For the acetic acid, only less than half of the calcium could be dissolved and the extraction kinetics were very slow, indicating that the extraction process only occurs at the surface of the wollastonite particles. Other studies have used oxalic acid, phosphoric acid, phthallic acid and combinations of these acids. However, although some of the above acids extracted the Mg2+ from serpentine, the obtained solutions do not precipitate magnesite, which may be due to the formation of complexes.
Consequently, mineral carbonation will only become a viable cost-effective sequestration technology when faster reaction routes under milder regimes in a continuous integrated process are developed.